Why we must use freshly prepared NaOH/SDS for miniprep - (Feb/06/2006 )

the 'precipitate' you are looking at is SDS. SDS does come out of solution when it gets cold. SDS precipitation is especially a problem during winter.

Sodium hydrogen carbonate and disodium carbonate are fully water soluble.

-perneseblue-

Yes it does. Look up Brodie's solution as applied to the old Warburg respirator by which expired carbon dioxide was measured as the precipitated salt. It depends on the concentration of NaOH. As for SDS - what concentration is used and why would expect precipitation failing saturation/lowered temp?

-jorge1907-

The typical NaOH/SDS solution has 1%w/v SDS and 0.2M NaOHIf you observed this solution when cold, you will see flakes of SDS precipitating. The appearance is quite dissimilar

Given the solubility of sodium hydrogen carbonate, the solution needs to react to alot of carbon dioxide to produce sufficient quantities of sodium hydrogen carbonate. Something that is hard to obtain in a closed bottle. I am quite sure you can rule out sodium carbonate precipitation.

I think there is a small mistake concerning your description of the Warburg respirator.

As best as I can understand, the Brodie solution and Warburg respirator are part of systems that use a manometer to measure partial pressure of a gas being emitted within a sample (usually CO2 or O2). As an example; CO2 in a gas sample reacts with a chemical (eg KOH), causing a reduction in volume of gas within the manometer. This results in a movement of a liquid of known density by a specific volume. Knowing the volume of liquid displaced and its density, you can then work out the partial pressure of the CO2 within the sample.

-perneseblue-

stick to what you know - and it's not the Warburg

-jorge1907-

stick to what you know - and it's not the Warburg. In fact - so should I. It's been alot off years and memory may betrray me - it may be the change in weight as the carbon dioxide is effectively captured rather than direct precipitation. I do dispute your solubility numbers as they speak top the absolute rahter than a more complex slution of higher pH. In any casde and at that caustic conc., I doubt my explanation works.

To your point - what's the solubility of SDS and how do you know the flakes are indeed SDS?

-jorge1907-

QUOTE (jorge1907 @ Aug 18 2008, 11:33 PM)

stick to what you know - and it's not the Warburg.

In fact - so should I. It's been alot off years and memory may betrray me - it may be the change in weight as the carbon dioxide is effectively captured rather than direct precipitation. I do dispute your solubility numbers as they speak top the absolute rahter than a more complex slution of higher pH. In any casde and at that caustic conc., I doubt my explanation works.

To your point - what's the solubility of SDS and how do you know the flakes are indeed SDS?

A more alkaline pH only means I should be using solubility values of disodium carbonate which is even more water soluble (My mistake). This situation is not complex. The solubility calculations for sodium carbonate/sodium hydrogen carbonate/Carbon dioxide system at a certain pH can be done by any A level chemistry student. It is an application solubility constants and partial pressure.

I know that these flakes are SDS, because I have made 10% SDS solutions and seen that solution precipitate during winter months. I have also seen crystals sodium carbonate that have precipitated out in solution. The structures are completely different.